View
42
Download
0
Category
Tags:
Preview:
DESCRIPTION
Properties of Carbon Element. Properties of Carbon Element. We have learnt carbon element as the basis of Organic Chemistry. Why does carbon can form so many different compounds? Carbon is found the second period group IVA of the periodic table. - PowerPoint PPT Presentation
Citation preview
Properties of Carbon Element
Properties of Carbon Element
• We have learnt carbon element as the basis of Organic Chemistry.
• Why does carbon can form so many different compounds?
• Carbon is found the second period group IVA of the periodic table.
EXCEPTIONS OF CARBON COMPUDS WHICH ARE NOT ORGANIC
• oxides of carbon (CO2, CO)
• carbonates,bicarbonates(NaHCO3,CaCO3)• cyanides (NaCN, etc)
Properties of Carbon Element
• The Lewis structure for carbon shows 4 unpaired valence electrons.
• To fulfill the octet rule, a carbon atom needs 4 more electrons.
• A carbon atom may form 4 covalent bonds and is capable of forming long chains with single, double or triple bonds between carbon atoms.
• These chains may be continuous (straight) or branched.
• The 2 ends of a chain can bond together to form a ring.
• Carbon compounds are divided into classes based on their chemical similarity.
Hydrocarbons• Hydrocarbons are compounds containing hydrogen and
carbon. Hydrocarbons may have different numbers of bonds between carbon atoms.
• The four hydrocarbon classes are: alkane (single bond), alkene, (double bond), alkyne (triple bond), aromatic (benzene ring).
• Alkanes contain only single C-C bonds. They contain as many hydrogen atoms as possible, and are said to be saturated.
• Hydrocarbons containing double or triple bonds are unsaturated.
• A homologous series is series of compounds that differ by a
constant increment. Aromatic hydrocarbons include a benzene ring- 6 carbon atoms with all the bonds alternating between a single and a double bond.
Properties of Carbon Element
Carbon is unique– It has 6 electrons in its outer
shell arranges 1s22s2sp2
– It has room for 4 bonds to 4 other atoms.
– Carbon-to-carbon bonds can be single (A),
– double (B), or – triple (C).
– Note that in each example,
each carbon atom has four dashes, which represent four bonding pairs of electrons, satisfying the octet rule.
Alkanes Alkenes Alkynes SATURATED means that each carbon is bonded to four other atoms through single covalent bonds. Hydrogen atoms usually occupy all available bonding positions after the carbons have bonded to each other.
UNSATURATED hydrocarbons contain either double or triple bonds. Since the compound is unsaturated with respect to hydrogen atoms, the extra electrons are shared between 2 carbon atoms forming double or triple bonds.
PARAFFINS which is derived from a Latin word meaning "little activity", and means that the compounds are very unreactive.
Alkenes are also called OLEFINS because they form oily liquids on reaction with chlorine gas.
Alkynes are also generally known as ACETYLENES from the first compound in the series.
HYDROCARBONS
HydrocarbonsHydrocarbons
C C C C
C C
C
C
C
C
C
C
H
H
H
H
H
H
C C C C C
H
H
H
H
H H
H
H
H
H
H
H
Alkanes Alkenes
Alkynes Aromatics
C C C C C
H
H
H
H
H
H
H
H
H
H
C C C C CH
H
H
H
H
H
H
H
Properties of Carbon Element
• A)The carbon atom forms bonds in a tetrahedral structure with a bond angle of 109.5O.
• (B) Carbon-to-carbon bond angles are 109.5O, so a chain of carbon atoms makes a zigzag pattern.
• (C) The unbranched chain of carbon atoms is usually simplified in a way that looks like a straight chain, but it is actually a zigzag, as shown in (B).
Properties of Carbon Element
Carbon-to-carbon chains can be
• (A) straight, • (B) branched, or • (C) in a closed ring.
• (Some carbon bonds are drawn longer, but are actually the same length.)
Why does carbon can form so many different compounds?
• There are now more than ten million organic compounds known by chemists.
• Many more undoubtedly exist in nature, and organic chemists are continually creating (synthesizing) new ones.
• Carbon is the only element that can form so many different compounds because each carbon atom can form four chemical bonds to other atoms, and because the carbon atom is just the right, small size to fit in comfortably as parts of very large molecules.
• Having the atomic number 6, every carbon atom has a total of six electrons.
• Two are in a completed inner shell, while the other four are valence electrons—outer electrons that are available for forming bonds with other atoms.
Why does carbon can form so many different compounds?
• The carbon atom's four valence electrons can be shared by other atoms that have electrons to share, thus forming covalent (shared-electron) bonds.
• They can even be shared by other carbon atoms, which in turn can share electrons with other carbon atoms and so on, forming long strings of carbon atoms, bonded to each other like links in a chain.
• Silicon (Si), another element in group 4A of the periodic table, also has four valence electrons and can make large molecules called silicones, but its atoms are too large to fit together into as great a variety of molecules as carbon atoms can.
Why does carbon can form so many different compounds?
• Carbon's ability to form long carbon-to-carbon chains is the first of five reasons that there can be so many different carbon compounds; a molecule that differs by even one atom is, of course, a molecule of a different compound.
• The second reason for carbon's astounding compound-forming ability is that carbon atoms can bind to each other not only in straight chains, but in complex branchings, like the branches of a tree.
• They can even join "head-to-tail" to make rings of carbon atoms. • There is practically no limit to the number or complexity of the branches
or the number of rings that can be attached to them, and hence no limit to the number of different molecules that can be formed.
• The third reason is that carbon atoms can share not only a single electron with another atom to form a single bond, but it can also share two or three electrons, forming a double or triple bond.
• This makes for a huge number of possible bond combinations at different places, making a huge number of different possible molecules.
• And a molecule that differs by even one atom or one bond position is a molecule of a different compound.
Why does carbon can form so many different compounds
• The fourth reason is that the same collection of atoms and bonds, but in a different geometrical arrangement within the molecule, makes a molecule with a different shape and hence different properties.
• These different molecules are called isomers.• The fifth reason is that all of the electrons that are not being
used to bond carbon atoms together into chains and rings can be used to form bonds with atoms of several other elements.
• The most common other element is hydrogen, which makes the family of compounds known as hydrocarbons.
• But nitrogen, oxygen, phosphorus, sulfur, halogens, and several other kinds of atoms can also be attached as part of an organic molecule.
• There is a huge number of ways in which they can be attached to the carbon-atom branches, and each variation makes a molecule of a different compound.
The Greater Stability of C-C Bonds
• Since the average bond dissociation energy of C-C is greater than the average bond energies between different atoms.
• Thus the energy released when carbon atom bonds to another carbon atom is greater than the energy released when the other atoms like B,N,O,Si,P and S bonds to each other.
• Thus C-C bond is more stable than the others like B-B,N-N,
O-O,Si-Si,P-P and S-S.
Bonding Atoms
Bond Energy (kJ mol-1)
B-B 293 C-C 343 N-N 163 O-O 157 Si-Si 222 P-P 201 S-S 266
Ability to Form Chains Between Their Atoms
• The atoms closer to C in the periodic table are B,N,O,Si,P and S.
• The ability of these atoms to bond each other to form chains is lower than C.
• For examle Si can produce chains made of at most 11 atoms of it and N at most three atoms it.
• Although the ability to form chains between their atoms for P and S is greater than Si and N but it is very much smaller compared to C.
Ability to Form Chains Between Their Atoms
The greater ability of carbon to form chains compared to atoms closer to it in the periodic table can be explained by two reasons:
1. The average bond dissociation energies of them is lower than that of carbon.
2. The electronegativity values B,Si and P lower than that of C.atoms.Thus the attraction forces between these atoms are smaller than that of carbon.This is also true when these atoms are bonded to the other atoms like hydrogen or halogens.
Li Be B C N O F 1,0 1,5 2,0 2,5 3,0 3,5 4,0 Na Mg Al Si P S Cl 0,9 1,3 1,6 1,9 2,2 2,5 3,2
Electronegativity values of some elements ording to Pauling’s Scale
Electronegativity
• Electronegativity:Electronegativity: – a measure of an atom’s attraction for the
electrons it shares with another atom in a chemical bond
• Pauling scalePauling scale– generally increases left to right in a row– generally increases bottom to top in a
column
Greater Bonding Capacity of C compared to N and O
The electronegativity values of N and O are greater than that of C. But their bonding capacities are smaller than that of C since they have lower number of unpaired electrons.
Lewis Dot Diagrams of Selected Elements
Summary…
• Compared to C atom B,Si,P,N and O atoms can not be expected to form greater number of compounds and unbrached and branched chains and cyclic compounds.
• Carbon compounds are more stable than Si4,P4,O3,S8 and B4 molecules.
Electron Configuration of Elements
Lewis Dot Structures…• Gilbert N. Lewis• Valence shell:Valence shell:
– the outermost occupied electron shell of an atom• Valence electrons:Valence electrons:
– electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions
• Lewis dot structure:Lewis dot structure: – the symbol of an element represents the nucleus and
all inner shell electrons– dots represent valence electrons
Lewis Dot Structures
• Table 1.4 Lewis Dot Structures for Elements 1-18
N OB
H
Li Be
Na
He
Cl
F
S
Ne
Ar
C
SiAl P
1A 2A 3A 4A 5A 6A 7A 8A
Mg ::
::
::
.
.
.
.
.
.
.
..
..
. .
.
.
.
:
:
:
::::::::
::::::.
:::
:
Lewis Model of Bonding…
• Atoms bond together so that each atom acquires an electron configuration the same as that of the noble gas nearest it in atomic number– an atom that gains electrons becomes an anionanion– an atom that loses electrons becomes a cationcation– the attraction of anions and cations leads to the
formation of ionic solidsionic solids– an atom may share electrons with one or more atoms
to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bondcovalent bond
– bonds may be partially ionic or partially covalent; these bonds are called polar covalent bondspolar covalent bonds
Covalent Bonds!• The simplest covalent bond is that in H2
– the single electrons from each atom combine to form an electron pair
– the shared pair functions in two ways
– simultaneously; it is shared by the two atoms and fills the valence shell of each atom
• The number of shared pairs
– one shared pair forms a single bond
– two shared pairs form a double bond
– three shared pairs form a triple bond
H H H-H+ • H0 = -435 kJ (-104 kcal)/mol•
Hydrogen Molecule Formation
0.74 A
- 436
0
H – H distance
En
erg
y (K
J/m
ol)
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318
no interaction
increasedattraction
balanced attraction& repulsion
increasedrepulsion
Potential Energy Diagram - Attraction vs. Repulsion
(internuclear distance)
Lewis Structures!• To write a Lewis structure
– determine the number of valence electrons
– determine the arrangement of atoms
– connect the atoms by single bonds
– arrange the remaining electrons so that each atom has a complete valence shell
– show a bonding pair of electrons as a single line
– show a nonbonding pair of electrons as a pair of dots
– in a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons
Table of Lewis Structures!
• In neutral molecules– hydrogen has one bond– carbon has 4 bonds and no lone pairs– nitrogen has 3 bonds and 1 lone pair– oxygen has 2 bonds and 2 lone pairs– halogens have 1 bond and 3 lone pairs
H2O (8) NH3 (8) CH4 (8) HCl (8)
C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)
H-O-H H-N-HH
H-C-HH
HH-Cl
H-C C-HH
HC O
H
HC C
H
HO O
CH H
O
Ethylene
Hydrogen chlorideMethaneAmmoniaWater
Carbonic acidFormaldehydeAcetylene
Resonance!• In chemistry, resonance is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. • A molecule or ion with such delocalized electrons is represented by several structures called resonance structures.
Tautomerization• Tautomerization usually involves the movement of a hydrogen atom between a different location on the molecule, resulting in two or more molecular structures.• These structures are called tautomers, which exist in dynamic equilibrium with each other.
Enol form Keto form
Molecular Geometry and Bonding Theories
molecular formula
structural formula
molecular shape
ball-and-stick model
CH4 C
H
H
HH
H
H
H
H
109.5o
C
tetrahedrontetrahedralshape ofmethane
CH
H
H
H
Methane & Carbon Tetrachloride
molecular formula
structural formula
molecular shape
ball-and-stick model
CH4 C
H
H
HH
H
H
H
H
109.5o
C
CCl4
space-filling model
C
Cl
Cl
ClCl
Molecular Geometry
Linear Trigonal planar
Tetrahedral
Trigonal pyramidalBent
107.3o104.5o
H2O CH4 AsCl3 AsF5 BeH2 BF3 CO2
180o
H
H
H
H
109.5o
C
C109.5o
H
HHH
N107o HH
H
..
O104.5o H
H
..
..
CH4, methane NH3, ammonia H2O, water
Molecular ShapesThree atoms (AB2) Four atoms (AB3)
Five atoms (AB4)
•Linear (180o)•Bent
•Trigonal planar (120o)•Trigonal pyramidal
•Tetrahedral (109.47o)
B BA
B
B
A
B
linear trigonal planar
B
A
BB
B
tetrahedral
Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 313.
Bonding and Shape of Molecules
Number of Bonds
Number of Unshared Pairs Shape Examples
2
3
4
3
2
0
0
0
1
2
Linear
Trigonal planar
Tetrahedral
Pyramidal
Bent
BeCl2
BF3
CH4, SiCl4
NH3, PCl3
H2O, H2S, SCl2
-Be-
B
C
N
:
O
:
:
CovalentStructure
Molecular Shapes
AB2
Linear
AB3
Trigonal planar
AB4
Tetrahedral
AB2EAngular or Bent
AB3ETrigonal
pyramidal
AB2E2
Angular or Bent
Molecular Polarity
Molecular Structure
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Dipole Moment
• Direction of the polar bond in a molecule.
• Arrow points toward the more electronegative atom.
H Cl+ -
Determining Molecular Polarity
• Depends on:– dipole moments– molecular shape
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Determining Molecular Polarity
• Nonpolar Molecules– Dipole moments are symmetrical and
cancel out.
BF3
F
F F
B
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Determining Molecular Polarity
• Polar Molecules– Dipole moments are asymmetrical and
don’t cancel .
netdipolemoment
H2OH H
O
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
CHCl3
H
Cl ClCl
Determining Molecular Polarity
• Therefore, polar molecules have...– asymmetrical shape (lone pairs) or – asymmetrical atoms
netdipolemoment
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Dipole Moment
Nonpolar
Polar
....
H H
O
C OO
Bond dipoles
Overall dipole moment = 0
Bond dipoles
Overall dipole moment
The overall dipole moment of a moleculeis the sum of its bond dipoles. In CO2 thebond dipoles are equal in magnitude butexactly opposite each other. The overall dipole moment is zero.
In H2O the bond dipoles are also equal inmagnitude but do not exactly oppose eachother. The molecule has a nonzero overall dipole moment.
221
dqqk
F Coulomb’s law = Q r Dipole moment,
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 315
......
Polar and Nonpolar Molecules
H Cl
Polar
A molecule has a zero dipole moment because their dipoles cancel one another.
H HO
PolarF F
B
F
Nonpolar
HH
H
N
Polar
Cl
ClC
Cl
Nonpolar Polar
Cl
HC
Cl
H
H
Be
H
H
BeH2
s p
Formation of BeH2 using pure s and p orbitals
The formation of BeH2 using hybridized orbitals
atomic orbitals atomic orbitals
Be
s p
Be H
H
s p
atomic orbitals
hybrid orbitals
No overlap = no bond!
sp p
Be HH
All hybridized bonds have equal strength and have orbitals with identical energies.
BeH2Be
Be = 1s22s2
Hybridization - The Blending of OrbitalsHybridization - The Blending of Orbitals
Poodle
+
+Cocker Spaniel
=
=
=
=
+
+s orbital p orbital
Cockapoo
sp orbital
sp Hybrid Orbitals
Ground-state Be atom
1s 2s 2p
1s 2s 2p
Be atom with one electron “promoted”
s
px py pz
sp
hybrid orbitals
En
erg
y
hybridize
s orbital p orbital
two sp hybrid orbitals sp hybrid orbitals shown together(large lobes only)
1s sp 2p
Be atom of BeH2 orbital diagram
H HBe
n = 1
n = 2
sp Animation
sp2 Hybrid Orbitals
2s 2p
Ground-state B atom
s
px py pzEne
rgy
sp2 2p
B atom of BH3 orbital diagram
hybridize
s orbital
2s 2p
B atom with one electron “promoted”
sp2
hybrid orbitals
p orbitals sp2 hybrid orbitals shown together
(large lobes only)three sps hybrid orbitals
H
H
HB
Sp2 Animation
s
px py pz
Carbon 1s22s22p2
Carbon could only make two bondsif no hybridization occurs. However,carbon can make four equivalent bonds.
sp3
hybrid orbitals
En
erg
y
sp3
C atom of CH4 orbital diagram
B
A
BB
B
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 321
sp3 Hybrid Orbitals
Sp3 Animation
Multiple Bonds
2s 2p 2s 2p sp2 2p
promote hybridize
C C
H
H H
H
C2H4, ethene
one bond and one bond
H
H
CC
H
H
H
H
CC
H
H
Two lobes ofone bond
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326
Multiple Bonds
2s 2p 2s 2p sp2 2p
promote hybridize
C C
H
H H
HC2H4, ethene (ethylene)
one bond and one bond
H
H
CC
H
H
H
H
CC
H
H
Two lobes ofone bond
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326
C C
H
H
sp2
sp2
sp2
H
H
sp2
sp2
sp2
p p
p p
Sigma and pi Bonds Animation
bond
Internuclear axis
p p
Two atomic p orbitals Pi (p) Bond
3D view of Pi (p) Bond
(pi) Bond – overlap of two p orbitals oriented perpendicular to the line connecting the nuclei.
bond
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 326
H s
Orbital Picture of Ethylene
C C
H s 1
sp2
sp2
sp2
sp2
sp2
sp2
H s
H s
px px
CH
HC
H
H
Ethylene Animation
Bonding in Formaldehyde
bonds in Benzene
H
CHC
H
C
H
C
H C
H
C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
C6H6 = benzene
2p atomic orbitals in Benzene
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
bonds and bonds
H
C
HCH
C
H
C
H C
H
C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
bonds in BenzeneH
CH
C
HC
H
C
H
C
H
C
HC
H
C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
bondsH
C
H
C
HC
H
C
H
C
H
C
HC
H
C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
Ethyne sp Hybridization
Ethyne sp Hybridization
Orbital Picture of Acetylene
C C
spsp
px
py
px
py
H s H ssp sp
1
2
C C HH
Acetylene Animation
Number of electron domains
Electron-domain geometry
Predicted bond angles
TetrahedralTrigonalplanar Tetrahedral
109.5o 120o 109.5o
C C OH H
H
H O
4 3 4
Acetic Acid, CH3COOH
Hybridization of central atom sp3 sp2
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 314
sp3
VSEPR Theory
• Valence Shell Electron Pair Repulsion Theory
• Electron pairs orient themselves in order to minimize repulsive forces.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
......
The Shapes of Some Simple ABn Molecules
H HB
Linear
O OS
BentF F
B
F
Trigonalplanar
FF
F
N
Trigonalpyramidal
SF6
H2O
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 305
VSEPR Theory
• Types of e- Pairs– Bonding pairs - form bonds– Lone pairs - nonbonding electrons
Lone pairs repel
more strongly than
bonding pairs!!!
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
VSEPR Theory
• Lone pairs reduce the bond angle between atoms.
Bond Angle
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
• Draw the Lewis Diagram.
• Tally up e- pairs on central atom.– double/triple bonds = ONE pair
• Shape is determined by the # of bonding pairs and lone pairs.
Know the 5 common shapes
& their bond angles!
Determining Molecular Shape
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Common Molecular Shapes
2 total
2 bond
0 lone
LINEAR180°BeH2
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
B BA
3 total
3 bond
0 lone
TRIGONAL PLANAR
120°
BF3
Common Molecular Shapes
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
B
B
A
B
4 total
4 bond
0 lone
TETRAHEDRAL
109.5°
CH4
Common Molecular Shapes
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
B
A
BB
B
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°
NH3
Common Molecular Shapes
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
4 total
2 bond
2 lone
BENT
104.5°H2O
Common Molecular Shapes
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
• CO2
O C O2 total
2 bond
0 lone LINEAR
180°
Examples
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
• PF3
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°
F P FF
Examples
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
No Loners Animation
Loners Animation
References
• http://science.jrank.org/pages/1202/Carbon-Why-carbon-special.html
• http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/perlewis.html
• www.nisd.net/communicationsarts/pages/chem
Recommended